Chemistry · Chemical Kinetics · NEET
Because most collisions are too weak or hit at the wrong angle. For a reaction, two things must both be true: (1) the colliding molecules must have energy equal to or greater than the threshold energy, so old bonds can break; and (2) they must collide in the correct orientation, so the right atoms line up to form new bonds. If either one is missing, the molecules just bounce off unchanged. Only the collisions that satisfy BOTH are called effective collisions, and only these lead to products.
Threshold energy is the minimum total energy the colliding molecules must have for a reaction to occur. Most molecules have less than this. Activation energy (Ea) is the EXTRA energy that reactant molecules must absorb to reach the threshold value. In short: Threshold energy = Average energy of reactants + Activation energy. A molecule below the threshold cannot react no matter how many times it collides.
No. Even a very energetic collision fails if the molecules are facing the wrong way. The reacting atoms must point toward each other so that the new bond can form. For example, if the wrong ends of two molecules meet, the collision is elastic and no product forms. This is why collision theory says a reaction needs enough energy AND proper orientation together. NEET often tests this — 'orientation' is a real, separate condition, not just extra detail.
Increasing concentration puts more molecules in the same volume, so molecules collide more often. This increases the collision frequency (number of collisions per unit time). It does NOT change threshold energy, activation energy, or the heat of reaction (ΔH) — those depend on the nature of the reactants and the energy barrier, not on how crowded the molecules are. This exact idea was asked in NEET 2020.
Collision frequency (often written Z) is the TOTAL number of collisions per second — most of these do nothing. Effective collisions are only the small fraction that have enough energy and the right orientation, so they actually form products. Rate depends on effective collisions, not just total collisions. That is why a reaction can be slow even when billions of collisions happen every second.
They are closely linked. The Arrhenius factor e^(−Ea/RT) in k = A·e^(−Ea/RT) represents the FRACTION of collisions that have energy ≥ threshold. The pre-exponential factor A is related to collision frequency and includes the orientation (steric) requirement through a probability factor P. So collision theory gives the physical meaning behind the Arrhenius equation.
An increase in the concentration of the reactants of a reaction leads to change in:
Activation energy of any chemical reaction can be calculated if one knows the value of:
Try the real previous-year questions from this chapter — each with the answer and a full solution.
(1) The colliding molecules must have energy equal to or above the threshold energy, and (2) they must collide with the proper orientation. Both must be satisfied for a reaction to occur.
Yes. Raising the temperature gives more molecules energy above the threshold value, so a larger fraction of collisions become effective. This is the main reason rate increases sharply with temperature.
P is a correction factor in collision theory that accounts for the orientation requirement. It shows that not every energetic collision has the right geometry, so the real rate is lower than the simple energy-based prediction.
It works well for simple bimolecular reactions of gases. For complex reactions and reactions in solution it is only approximate, because it treats molecules as hard spheres and ignores their internal structure.
NEET regularly asks which quantities change with concentration or temperature, and why energy plus orientation are both needed. Knowing effective collisions vs collision frequency helps you avoid common traps in Chemical Kinetics.
